Electrolysis — Railway Group D Study Notes

Overview

Electrolysis is the chemical decomposition of a substance using direct electric current. This topic is critical for Railway Group D as questions frequently test the basic principles of electrolytic cells, practical applications like electroplating, and Faraday's laws of electrolysis. Unlike galvanic cells (batteries) that produce electricity from spontaneous reactions, electrolytic cells use electrical energy to drive non-spontaneous chemical reactions.

Understanding electrolysis requires grasping three core elements: how the electrolytic cell works, what Faraday's laws quantify, and real-world applications such as electroplating and purification of metals. Exam questions typically ask about electrode reactions (anode vs cathode), the relationship between charge passed and mass deposited, and practical uses in industry. Master the direction of electron flow, ion movement, and the calculations involving Faraday's constants to tackle quantitative problems confidently.

Expect 1–3 questions from this topic, often integrated with electrochemistry or practical chemistry applications. Focus on understanding the mechanism rather than rote memorization—this helps answer both theoretical and numerical questions effectively.

Key Concepts

• **Electrolytic cell structure**: Consists of two electrodes (anode and cathode) immersed in an electrolyte solution or molten ionic compound, connected to an external battery or DC source. The positive terminal connects to the anode, negative to the cathode.

• **Electrode identification**: In electrolysis, the anode is positive (oxidation occurs—loss of electrons), and the cathode is negative (reduction occurs—gain of electrons). Remember: "An Ox, Red Cat" (Anode Oxidation, Reduction Cathode).

• **Ion migration**: Cations (positive ions) move toward the cathode, while anions (negative ions) move toward the anode. Discharge occurs at electrodes based on preferential discharge series—less reactive ions discharge first.

• **Electrolyte**: Can be aqueous solution (like copper sulphate, acidified water) or molten salts (like sodium chloride). Water itself can undergo electrolysis to produce hydrogen and oxygen gases.

• **Faraday's first law**: The mass of substance deposited or liberated at an electrode is directly proportional to the quantity of electricity (charge) passed through the electrolyte.

• **Faraday's second law**: When the same quantity of electricity passes through different electrolytes, the masses of substances deposited are proportional to their equivalent weights (atomic weight/valency).

• **Electroplating**: Practical application where a thin layer of one metal is deposited onto another using electrolysis to prevent corrosion, improve appearance, or reduce cost. The object to be plated acts as the cathode.

• **Faraday constant**: One Faraday (F) equals 96,500 coulombs—the charge carried by one mole of electrons. This links chemical change to electrical charge quantitatively.

Formulas / Key Facts

**Faraday's First Law Formula:** m = Z × Q or m = Z × I × t Where m = mass deposited (grams), Z = electrochemical equivalent (g/C), Q = charge (coulombs), I = current (amperes), t = time (seconds)

**Electrochemical Equivalent (Z):** Z = E/F = (Atomic weight)/(Valency × 96,500) Where E = equivalent weight, F = Faraday's constant (96,500 C/mol)

**Faraday's Second Law:** m₁/m₂ = E₁/E₂ Masses deposited by same charge are in the ratio of their equivalent weights.

**Charge-Mole Relationship:** Q = n × F Where n = number of moles of electrons, F = 96,500 C

**Key Fact**: 1 Faraday deposits 1 gram equivalent weight of any substance.

**Electrolysis of Water:** 2H₂O → 2H₂ + O₂ (at cathode: hydrogen; at anode: oxygen in 2:1 volume ratio)

**Preferential Discharge**: Metals below hydrogen in reactivity series (Cu, Ag, Au) discharge at cathode in preference to hydrogen ions.

**Electroplating Requirements**: Electrolyte contains salt of plating metal; anode is made of plating metal (replenishes ions); cathode is the object to be plated.

Worked Examples

**Example 1: Mass Deposited** *Question*: A current of 2 amperes is passed through copper sulphate solution for 16 minutes 5 seconds. If the electrochemical equivalent of copper is 0.000329 g/C, find the mass of copper deposited.

*Solution*: Step 1: Convert time to seconds: 16 min 5 sec = (16 × 60) + 5 = 965 seconds Step 2: Calculate charge: Q = I × t = 2 × 965 = 1930 coulombs Step 3: Apply Faraday's first law: m = Z × Q = 0.000329 × 1930 = 0.635 grams **Answer**: 0.635 grams of copper deposited

**Example 2: Using Faraday Constant** *Question*: How many grams of silver (atomic weight 108, valency 1) will be deposited by passing 0.5 Faraday of electricity?

*Solution*: Step 1: 1 Faraday deposits 1 equivalent weight Step 2: Equivalent weight of silver = 108/1 = 108 grams Step 3: 0.5 Faraday deposits = 0.5 × 108 = 54 grams **Answer**: 54 grams of silver

**Example 3: Comparative Deposition** *Question*: If 3.15 g of copper is deposited by a certain quantity of electricity, how much aluminum (atomic weight 27, valency 3) will be deposited by the same charge? (Copper: atomic weight 63.5, valency 2)

*Solution*: Step 1: Find equivalent weights: Cu = 63.5/2 = 31.75; Al = 27/3 = 9 Step 2: Apply Faraday's second law: m(Al)/m(Cu) = E(Al)/E(Cu) Step 3: m(Al)/3.15 = 9/31.75 Step 4: m(Al) = 3.15 × (9/31.75) = 0.893 grams **Answer**: 0.893 grams of aluminum

Common Mistakes

**Mistake 1**: Confusing anode and cathode polarity in electrolytic vs galvanic cells. → **Fix**: In electrolysis, anode is POSITIVE (connected to + terminal), cathode is NEGATIVE. Oxidation always at anode, reduction always at cathode regardless of cell type.

**Mistake 2**: Using atomic weight directly in Faraday's law calculations instead of equivalent weight. → **Fix**: Always divide atomic weight by valency to get equivalent weight. For Cu²⁺, use 63.5/2 = 31.75, not 63.5.

**Mistake 3**: Wrong unit conversion—forgetting to convert minutes to seconds or milliamperes to amperes in Q = I × t. → **Fix**: Standardize units before calculation: time in seconds, current in amperes, charge in coulombs.

**Mistake 4**: Thinking the object to be electroplated should be the anode. → **Fix**: The object receiving the coating must be the cathode (negative electrode) where metal cations deposit. The anode is made of the plating metal.

**Mistake 5**: Assuming water electrolysis always produces H₂ and O₂ in all solutions. → **Fix**: Preferential discharge applies—in CuSO₄ solution, Cu²⁺ discharges at cathode instead of H⁺ because copper is less reactive than hydrogen.

Quick Reference

• **Electrolysis = chemical decomposition using electric current; requires external DC source** • **Anode (+ve) → Oxidation; Cathode (−ve) → Reduction | Anions to anode, cations to cathode** • **Faraday's 1st Law: m = Z × I × t | Mass ∝ Charge passed** • **1 Faraday = 96,500 coulombs = charge on 1 mole electrons = deposits 1 gram equivalent** • **Electrochemical equivalent Z = (Atomic weight)/(Valency × 96,500)** • **Electroplating: Cathode = object to be plated; Anode = plating metal; Electrolyte = salt of plating metal**